7.4 Formal Charges and Resonance

Formal charges and resonance are crucial concepts in understanding molecular structure and stability. They help explain how electrons are distributed in molecules and why certain structures are more stable than others.

Calculating formal charges allows us to predict the most stable Lewis structures. Resonance forms show how electrons can be delocalized across molecules, leading to enhanced stability and unique chemical properties.

Formal Charges and Resonance

Formal charges in Lewis structures

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• Formal charge represents the difference between the number of valence electrons in an isolated atom and the number of electrons assigned to that atom in a Lewis structure
  • Valence electrons are the electrons in the outermost shell of an atom (Group 1A: 1, Group 2A: 2, Group 3A-8A: 3-8)
  • Assigned electrons include electrons in lone pairs and half of the electrons in bonding pairs
• Calculate formal charge using the formula: $Formal\,charge = Valence\,electrons - (Lone\,pair\,electrons + \frac{1}{2}Bonding\,pair\,electrons)$
  • Lone pair electrons are non-bonding electrons assigned to a specific atom (H2O, NH3)
  • Bonding pair electrons are electrons shared between atoms in a covalent bond (H2, Cl2)
• Carbon monoxide (CO) example:
  • Carbon has 4 valence electrons, 2 lone pair electrons, and 2 bonding pair electrons
    • $Formal\,charge\,(C) = 4 - (2 + \frac{1}{2}(2)) = +1$
  • Oxygen has 6 valence electrons, 4 lone pair electrons, and 2 bonding pair electrons
    • $Formal\,charge\,(O) = 6 - (4 + \frac{1}{2}(2)) = -1$

Evaluation with formal charges

• The most stable Lewis structure minimizes the sum of the absolute values of formal charges
  • Structures with formal charges closer to zero are generally more stable (CH4, H2O)
• If multiple structures have the same sum of absolute formal charges, the most stable structure will have:
  • Negative formal charges on more electronegative atoms (O, N, F)
  • Positive formal charges on less electronegative atoms (C, P, S)
• Comparing two Lewis structures for CO2 example:
  • Structure 1: C=O=O
    • $Formal\,charge\,(C) = 4 - (0 + \frac{1}{2}(8)) = 0$
    • $Formal\,charge\,(O) = 6 - (4 + \frac{1}{2}(4)) = 0$
    • Sum of absolute formal charges = 0
  • Structure 2: C-O=O+
    • $Formal\,charge\,(C) = 4 - (2 + \frac{1}{2}(4)) = 0$
    • $Formal\,charge\,(O^-) = 6 - (6 + \frac{1}{2}(2)) = -1$
    • $Formal\,charge\,(O^+) = 6 - (2 + \frac{1}{2}(4)) = +1$
    • Sum of absolute formal charges = 2
  • Structure 1 is more stable due to lower sum of absolute formal charges

Resonance forms of molecules

• Resonance occurs when a molecule can be represented by two or more Lewis structures
  • The actual structure is a hybrid of these resonance forms (benzene, nitrate ion)
• Resonance forms differ in the arrangement of electrons, but not in the positions of atoms
• Draw resonance forms using these steps:
  1. Start with one valid Lewis structure
  2. Move electrons (as lone pairs or bonding pairs) to create another valid Lewis structure
  • Do not change the positions of atoms
  • Maintain the octet rule for each atom
• Ozone (O3) example:
  • Resonance form 1: O=O-O
    • Central oxygen has a positive formal charge, terminal oxygen has a negative formal charge
  • Resonance form 2: O-O=O
    • Central oxygen has a positive formal charge, other terminal oxygen has a negative formal charge
  • The actual structure is a hybrid of these two forms, with a bond order of 1.5 between each oxygen atom
• Resonance involves electron delocalization, where electrons are spread over multiple atoms or bonds

Advanced Concepts in Resonance

• Molecular orbital theory provides a more accurate description of electron behavior in resonance structures
• Pi bonds play a crucial role in resonance, as they can be easily delocalized across multiple atoms
• Conjugation occurs when alternating single and double bonds allow for extended electron delocalization
• Aromaticity is a special case of resonance stability in cyclic, planar molecules with delocalized electrons